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• Dalton’s model proposed that the atom was the smallest particle of matter and was indivisible (like billiard balls) all matter was composed of atoms, the atoms of each element were different •Thompson’s Model: The positive charges fills the atom while the electrons were embedded throughout the atom (raisin bun model) •Thomson discovered the electron and since the electron was negative, but atoms neutral, there had to positive charges in atoms The positive fluid fills the atom while the electrons were embedded throughout the atom (raisin bun model) Thomson used a beam of cathode rays in a CRT with electric field and a magnetic field perpendicular to the direction of beam travel With only the electric field on, the beam was deflected • With only the magnetic field on, the cathode rays were deflected into a curved path When both fields were on, and the field strengths equal, the cathode rays were not deflected Thomson showed that they had a negative charge and that they had mass, so had to be particles coming from the atom Thomson determined the charge to mass ratio of the cathode rays using the curvature of the beam caused by the magnetic field Fc Fm 2 mv qvB r mv qB r v q B r m Thomson's experiment could not determine the mass or the charge, only the ratio Thompson determined the speed of the cathode rays: Since the beam was not deflected when both fields were adjusted correctly, Fe E Fe E q q Fm qvB Fe = Fm (balanced forces, 1st Law) E q qvB E v B Rutherford Model •Around 1911 Rutherford, Marsden and Geiger performed experiments to test the Thomson model •Alpha particles from radioactive sources were directed at thin gold foils (about 100 atoms thick) The Thomson model predicted that most of the alpha particles would go straight through, and only a few would be deflected at small angles since the electrons in the atom have much less mass than -particles. Rutherford experiment animation •Most of the -particles went straight through undeflected, some were deflected at angles of more than 10o and a few were deflected almost straight back •He concluded that most of the atom was empty space with most of the mass and all of the positive charge concentrated in a very small region (the nucleus) •Scattering angles indicated the size of the nucleus was about 10-15 to 10-14 m in radius Rutherford's model had 2 problems: 1) The model could not explain the light emitted from hot gases 2) Maxwell’s theory predicted that accelerating electrons would radiate EMR, so the electrons in Rutherford’s model should radiate a continuous spectrum and lose energy as they spiraled into the nucleus •Niels Bohr developed a theory to explain why the atom was stable: •He suggested that the electron could not lose energy continuously, but had to do so in “quantum jumps” and that the electrons moved around the nucleus in circular orbits, and only certain orbits were allowed •Each orbit represented a certain amount of energy and that the electron would move in the orbit without radiating energy (although this violated classical physics) The allowed orbits were called stationary states, a photon would be absorbed when the electron jumped from one state to a higher state, a photon would be emitted if the electron dropped to a lower state Bohr E-level applet The radius of each orbit was given by the equation r n r1 2 radius of Bohr orbit number of the orbit or energy level: 1, 2, 3, 4, etc. r1 = radius of 1st orbit, the ground state (5.29 x 10-11 m for hydrogen atom) (on data sheet under “Atomic Physics”) • Each higher orbit represents more energy. • The energy of the electron in each orbit was found by (on data sheet) Energy of electron in energy level n 1 En 2 E1 n Energy level Energy of electron in energy level 1 For hydrogen, E1 = -13.6 eV or – 2.18 x 10-18 J •The zero point for the energy is when the e- is an infinite distance from the nucleus •The negative sign means that the electron needs that much energy to be removed from the atom (to ionize the atom) •The energy of the photon that an atom absorbs or emits can be found by hf = Ef – Ei (energy in the final orbit – energy in the initial orbit) •When the atom is cool or unexcited, the electron is in the ground state (n = 1) •In a sample of excited hydrogen atoms, each electron will release a photon as it drops to a lower level, forming a bright line spectrum Where does this energy come from? EMR - photons • If the photon has exactly the right amount of energy, the atom will absorb this energy and the photon will no longer exist. High speed electrons (a la Frank Hertz exp) • If the electron has enough energy, the atom will absorb (in a collision) this energy and the remainder would exist as the kinetic energy of the electron Examples 1.Determine the wavelength of the fourth line in the Balmer series (nf = 2, ni = 6). = 4.09 x 10-7 m 2. If an electron absorbs exactly 2.86 eV from a photon, what was the electron’s stationary state, and where did it move to? When it moved back to its stationary state, what colour was the EMR emitted? 3. Determine the energy emitted by an electron as it drops from the 5th to the 3rd energy level of a hydrogen atom. E = -0.967 eV The atom lost 0.967 eV of energy. Summary •The Bohr model accurately explained atomic spectra (emission and absorption) •The energy of each orbit is quantized •Explained why atoms are stable •Explained the chemical and physical properties of the elements •Could not explain why electrons could only be found in certain orbits •Could not explain why some lines in the emission spectra were brighter than others •Worked only for hydrogen and atoms with a single electron (for example He+)